Chapter 13 Review
Vocabulary
Match each word to the sentence where it best
fits.
Section 13.1
acid
base
hydronium ion
Bronsted-Lowry
neutral
amphoteric
strong acid
weak base
weak acid
strong base
1. HCl is an example of a strong __________
because it releases hydrogen ions.
2. A solution that is _____________ contains more
hydroxide ions than hydrogen ions.
3. A solution that has a pH of seven is considered to
be ______________.
4. Water can accept an H+ ion or donate an H+ ion
and this makes it _____________.
5. A _____________ only ionizes partially in water
and yields fewer hydrogen ions than a
_____________ which ionizes 100% in water.
6. Ammonia, NH3 is an example of a ____________.
7. The __________________ definition of an acid
and a base explains which substance is a hydrogen
ion acceptor and which is a hydrogen ion donor.
11. Phenolpthalein is and example of an __________,
which is a substance that turns a different color as
the pH changes.
12. In the number 2.3 x 10-5, the negative five
represents an ____________. It is negative
because the number represents a value less than
one.
13. The equilibrium constant for water at 25oC and 1
atm is called the ________________ which
represents a value called Kw.
Section 13.4
neutralization
equivalence point
titration
salt
common ion
buffer
buffer capacity
14. A _____________ is an important laboratory
method used to determine the concentration of a
solution of an acid or base.
15. During a ______________ reaction, the acid and
the base react until the H+ ions = the OH- ions.
16. At the _________________ we can make the
important assumption that moles base = moles of
acid.
17. A solution that is able to resist small changes in
pH is called a _____________.
8. When water bonds to a hydrogen ion in an
aqueous solution the_ ____________ is formed in
solution
18. Potassium chloride, KCl is a _________ that is
formed from the positive ion of a base and the
negative ion of an acid.
9. Sodium hydroxide, NaOH is an example of a
_________________.
19. Buffers work based on the ______________
principle, which recognizes that an ion is produced
by two or more chemicals in the same solution.
Sections 13.2 and 13.3
exponent
logarithm
indicatorl
ion product constant
20. The ______________ tells us how much acid or
base a solution can neutralize before changing pH.
10. The __________ of 10,000 is 104.
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A NATURAL APPROACH TO CHEMISTRY
Conceptual Questions
32. Look at 30 (a ) and (b) above and explain water’s
role in each equation.
Section 13.1
21. Compare and contrast an acid and a base. Clearly
explain their chemical differences and give an
example of each.
22. Give two examples (each) of acidic and basic
substances that you use regularly in your daily
lives.
23. Explain what a “neutral” solution means in terms
of the chemical ions present.
24. Write a chemical equation that represents :
a) How a strong base ionizes in water.
b) How a strong acid ionizes in water.
25. What type of solution is formed when Group I and
II metals dissolve in water ? Explain and support
your answer using a chemical equation.
26. The hydrogen ion can be thought of as a “naked
proton.” Explain why this is important in terms of
chemical reactions.
27. Explain the main difference between the early
Arrhenius theory of acids and base and the
Bronsted-Lowry theory of acids and bases.
28. Explain how ammonia, NH3, acts as a BronstedLowry base when dissolved in water. Show the
chemical equation that represents this process.
33. Identify the following substances as strong acids
or strong bases
a) KOH
b) HNO3
c) Ba(OH)2
d) H2SO4
e) HCl
34. Write a chemical equation that shows how nitric
acid, HNO3, ionizes in anaqueous solution.
Section 13.2
35. Briefly explain what pH measures in an aqueous
solution.
36. Make a sketch of the pH scale and label the
numbers zero to 14. Label where on the scale
these substances would be located.
a) Lemon juice
b) baking soda
c) dish soap
d) soda water
e) blood
f) milk
37. The pH of a solution is 6.8. From this information
can you conclude that the solution is acidic?
Explain.
38. Write the equation that relates pH to pOH.
29. What ion is formed when an acid is dissolved in
water? Show the Lewis structure for this ion.
39. If you have a “strong” acid does that mean the pH
of the solution of this acid will be low? Explain.
30. In each of the following chemical reactions
identify the substance acting as the acid and the
substance acting as the base. Label (A)acid or
(B)base.
F- + H3O+
a) HF + H2O
40. Based on the pH scale sustances that are acidic
have a pH less than 7. What does this tell you
about the amount of [H+] relative to [OH-] in the
solution? Explain.
-
b) HNO2 + H2O
c) NH3 + H2O
+
NO2 + H3O
NH4+ + OH-
d)HC2H3O2 + NH3
C2H3O2- + NH4+
31. For number 30 (b) above clearly explain answer.
A NATURAL APPROACH TO CHEMISTRY
41. Explain how an aqueous solution can have a pH of
7. What would have to be true?
42. If the pH changes by 2 units by what factor does
the [H+] change?
43. Name two different types of indicators.
44. How are indicators used to help determine the pH
of a solution? Explain.
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Chapter 13 Review.
45. Sometimes the pH of a an acidic solution can be
zero or negative. Explain why.
Section 13.3
46. a)Write the chemical equation for the
autoionization of water.
b) The equilibrium constant, Kw, for this reaction
is 1.0 x 10-14 . Explain what the magnitude of the
Kw tells us about the reaction in the forward
direction.
56. a) Show the chemical reaction for the addition of
the salt, potassium nitrate KNO3, to water.
b) Will this aqueous solution conduct electricity ?
Explain.
57. List two common salts formed during the
neutralization of a strong acid and strong base.
58. Explain how adding 4.0 g of the salt sodium
acetate, NaC2H3O2, to an acidic solution would
affect the pH?
47. a) Write the chemical reaction for the strong acid,
hydrochloric acid HCl, ionizing in water.
b) Write the chemical reaction for the weak acid,
acetic acid HC2H3O2, ionizing in water.
c) Compare the weak and strong acid’s behavior in
water.
59. Predict how the overall pH would be affected, if
some sodium chloride, NaCl solid where added to
a neutral solution with a pH of 7.
48. Give an example of a weak base and a weak acid.
Quantitative Problems
49. Calculate the number of moles of NaOH that are
in 8.50 mL of a 0.450 M solution.
50. Predict the products for the following weak acid
and base ionizations shown below
a) HF + H2O
b) CH3NH2
c) HNO2 + H2O
51. If you have a 1.0 M soltution of HCl and a 1.0M
solution of HC2H3O2, List some things that will
be different in each of these solutions. If you did
not know which was which, how could you
determine which one was HCl?
Section 13.4
52. Give an example of the following acid - base
reactions, include reactants and products.
a) neutralization
b) metal and acid
53. What are the products of a neutralization reaction
between HCl and NaOH?
54. Do you think neutralization reactions always have
a pH of 7? Why or why not?
55. In the laboratory baking soda is often kept handy
to neutralize the any strong acid spills. How does
baking soda neutralize strong acids?
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60. a)What is unique about a buffer?
b) Can acetic acid act as a buffer? Explain.
Section 13.2
61. If a solution has an [H+] of 3.5 x 10-4, what is the
pH of the solution?
62. A solution is measured to have a pH of 5.60,
calculate the hydrogen ion concentration of the
solution?
63. Calculate the pH of the following solutions :
a) A glass of lemonade with a 4.30 x 10-5 M H+
concentration.
b) Salad dressing with an [H+] = 2.51 x 10-6
c) A glass of tomato juice with an [H+] = 6.40 x
10-5
d) A glass of milk with an [H+] = 7.94 x 10-9
64. Calculate the pH for each of the hydrogen ion
concentrations. State whether the pH is acidic,
basic or neutral.
a) [H+] = 2.45 x 10-4
b) [H+] = 2.11 x 10-11
c) [H+] = 4.96 x 10-7
d) [H+] = 1.00 x 10-7
e) [H+] = 3.51 x 10-14
65. Calculate the pH for each of the hydroxide ion
concentrations. State whether the pH is acidic,
basic or neutral.
a) [OH-] = 4.03 x 10-2
A NATURAL APPROACH TO CHEMISTRY
b) [OH-] = 9.55 x 10-9
c) [OH-] = 5.20 x 10-5
d) [OH-] = 7.26 x 10-12
Section 13.3
66. Calculate the [OH-] concentration for each of the
following [H+] concentrations. State whether the
pH is acidic, basic, neutral.
a) [H+] = 3.15 x 10-4
b) [H+] = 2.11 x 10-10
c) [H+] = 5.96 x 10-7
67. Calculate the [H+] concentration for each of the
following [OH-] concentrations. State whether the
pH is acidic, basic, neutral.
a) [OH-] = 3.23 x 10-2
b) [OH-] =8.65 x 10-8
c) [OH-] = 5.41 x 10-6
68. Calculate the hydrogen ion concentration and the
pH of each of the solutions of strong acids below.
a) 2.4 x 10-5 M HCl
b) 0.0045 M H2SO4
c) 5.51 x 10-4 HNO3
d) 6.41 x 10-4 HCl
69. Determine the pH of a 0.50 M solution of acetic
acid, HC2H3O2. Use a RICE table to do your
calculation. Ka for acetic acid = 1.8 x 10-5
70. Calculate the pH of a 0.10 M solution of lactic
acid. Lactic acid HC3H5O3 is a weak monoprotic
acid and it is the acid that sometimes makes your
muscles sore after you exercise. Ka = 1.4 x 10-4
71. Calculate the pH of a 0.25 M solution of ammonia,
NH3 a weak base. Use a RICE table to do your
calculation. Kb for ammonia = 1.8 x 10-5
72. Calculate the pH of a 0.085 M solution of
ethylamine, C2H5NH2. The Kb for ethylamine is
6.4 x 10-4.
73. Calculate the percent ionization for the acetic acid
in number 60) above. Show your work.
74. Codeine, C18H21NO3, is a weak organic base.
Calculate the Kb of a codeine solution with an
A NATURAL APPROACH TO CHEMISTRY
initial molarity of 4.80 x 10-3 M, given that this
codeine solution has a pH of 9.90.
Section 13.4
75. Calculate the molarity of an NaOH solution, if
26.0 mL of the solution are needed to neutralize
16.8 mL of a 0.298 M HCl solution. Remember
that “neutralize” means that the moles [H+] =
moles [OH-].
76. Calculate the volume in milliliters of a 1.26 M
NaOH solution required to titrate the solutions
shown below:
a) 25.00 mL of 2.35 M HCl solution.
b) 25.00 mL of 4.30 M HNO3 solution
c) 30.00 mL of 3.35 M HCl solution.
77. Calculate the volume of 0.80 M HCl solution
needed to neutralize each of the following basic
solutions.
a) 12.00 mL of a 0.400 M NaOH solution
b) 12.00 mL of a 0.200 M Ba(OH)2 solution
78. A 2.653 g sample of a solid monoprotic acid was
dissolved in 25.0 mL of distilled water. The
solution required 18.30 mL of a 0.160 M NaOH
solution to titrate. Calculate the molar mass of the
acid.
79. Acetic acid, CH3COOH is the acid present in
apple cider vinegar. A sample of 25.00 mL of
vinegar is titrated using a 1.10 M solution of
NaOH. This sample required 2.35 mL of NaOH to
neutralize it. Determine the concentration of the
acetic acid in the vinegar?
80. What volume of 0.152 M HCl is required to
neutralize 3.40 g of Mg(OH)2? Write out the
balanced chemical reaction that occurs to help
with this calculation.
81. How many milliliters of 0.120 M HCl are needed
to neutralize completely 31.00 mL of 0.150M
Ba(OH)2 solution? Write out the balanced
chemical reaction that occurs as part of your
answer.
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